4.2 Covalent bonding and Lewis structures
3 min read•july 30, 2024
Covalent bonds form when atoms share electrons, creating stable molecules. This sharing leads to specific bond properties, like length and angle, which influence molecular structure and behavior.
Lewis structures visually represent electron arrangements in molecules, helping predict their shape and properties. Understanding these concepts is crucial for grasping how atoms interact and form compounds.
Covalent bond formation and characteristics
Covalent bond formation
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- Covalent bonds form when atoms share one or more pairs of electrons
- Typically occurs between nonmetals (carbon, oxygen, nitrogen)
- Shared electrons are attracted to the nuclei of both atoms
- Attraction results in a stable bond between the atoms
- Number of shared electron pairs determines bond order (single, double, triple)
Covalent bond properties
- Covalent bonds are directional with specific bond lengths and angles
- is the distance between the nuclei of the bonded atoms
- Bond angle is the angle formed by the imaginary lines connecting the nuclei
- Strength of a covalent bond depends on bond order and electronegativity difference
- Higher bond order (double or triple bonds) results in stronger bonds
- Greater electronegativity difference leads to more polar covalent bonds
Lewis structures for molecules and ions
Constructing Lewis structures
- Lewis structures represent the arrangement of atoms and electrons in a molecule or polyatomic ion
- Each atom is represented by its chemical symbol
- are shown as dots surrounding the symbol
- Atoms are connected by lines representing shared electron pairs (covalent bonds)
- states that atoms tend to achieve a stable electron configuration with eight valence electrons (except hydrogen, which requires two)
- Lone pairs are unshared electron pairs not involved in bonding
Guidelines for drawing Lewis structures
- Place the least electronegative atom at the center (usually carbon or the least electronegative nonmetal)
- Arrange the remaining atoms around the central atom
- Distribute the available valence electrons as lone pairs or bonding pairs
- Assign formal charges to atoms to determine the most stable arrangement
- Formal charge = [# of valence electrons in free atom] - [# of non-bonding electrons] - (1/2)[# of bonding electrons]
- Aim for the most stable structure with the lowest formal charges and satisfying the octet rule
Molecular geometry and polarity
VSEPR theory and molecular geometry
- VSEPR (Valence Shell Electron Pair Repulsion) theory predicts molecular geometry based on electron domains (bonding and nonbonding) around the central atom
- Electron domains repel each other and arrange to minimize repulsion
- Five basic electron domain geometries: (2 domains), (3 domains), (4 domains), trigonal bipyramidal (5 domains), octahedral (6 domains)
- Molecular geometry is determined by the arrangement of atoms, not considering lone pairs
- Lone pairs affect bond angles and overall molecular shape (bent, trigonal pyramidal, seesaw, T-shaped, square pyramidal)
Molecular polarity
- Polarity depends on geometry and electronegativity differences between bonded atoms
- Polar molecules have an uneven charge distribution due to polar covalent bonds and asymmetric geometry (water, ammonia)
- Nonpolar molecules have a symmetric charge distribution, either due to absence of polar covalent bonds (O2, N2) or symmetric arrangement of polar bonds (CO2, BF3)
- Dipole moment is a vector quantity that measures the magnitude and direction of polarity
- Molecules with a net dipole moment are polar, while those with a zero net dipole moment are nonpolar
Resonance and molecular stability
Resonance structures
- Resonance occurs when a molecule or polyatomic ion can be represented by multiple valid Lewis structures
- Actual structure is a hybrid of these , with electron density distributed among them
- Resonance structures are not different molecules, but different representations of the same molecule or ion
- Resonance structures are connected by double-headed arrows, indicating they contribute to the overall electronic structure
Resonance and stability
- More resonance structures lead to greater stability due to electron delocalization
- Delocalization lowers the overall energy of the molecule or ion compared to any single resonance structure
- Stability gained through resonance is called resonance energy
- Resonance affects bond lengths, with actual bond lengths intermediate between single and double bonds
- Examples of molecules with significant resonance: benzene, nitrate ion, carbonate ion
Key Terms to Review (16)
Bond dissociation energy: Bond dissociation energy is the amount of energy required to break a specific bond in a molecule, resulting in the formation of neutral atoms or radicals. It reflects the strength of a chemical bond; the higher the bond dissociation energy, the stronger the bond. This concept is crucial in understanding both covalent bonding and the behavior of electrons in molecular structures, and it plays a significant role in predicting reaction pathways and molecular stability.
Bond Length: Bond length is defined as the average distance between the nuclei of two bonded atoms in a molecule. It plays a crucial role in understanding molecular geometry, stability, and reactivity, as well as the vibrational characteristics that can be observed through various spectroscopic techniques. The bond length is influenced by factors such as atomic size, bond order, and the presence of electron repulsion, all of which help determine how molecules interact with each other.
Carbon dioxide (CO2): Carbon dioxide (CO2) is a colorless, odorless gas that is essential for life on Earth, produced by the respiration of animals and the combustion of fossil fuels. It plays a crucial role in the greenhouse effect, where it traps heat in the atmosphere, contributing to global warming. Its molecular structure and bonding characteristics are key to understanding its behavior in various chemical reactions and its interactions with biological systems.
Delocalized electrons: Delocalized electrons are electrons that are not confined to a single atom or a single covalent bond but are spread over several atoms, contributing to the stability of a molecule. This behavior occurs in certain molecular structures, such as aromatic compounds and metals, where electrons can move freely across multiple nuclei. Delocalization helps explain properties like electrical conductivity and resonance stability.
Hybridization: Hybridization is the process of combining atomic orbitals to form new hybrid orbitals, which can then be used to form covalent bonds in molecules. This concept helps explain the geometry and bonding properties of molecules by accounting for the observed bond angles and shapes that cannot be explained by simple atomic orbitals alone. By mixing different types of orbitals, such as s, p, and sometimes d orbitals, hybridization allows for a more accurate representation of molecular structure.
Linear: In chemistry, the term 'linear' refers to a molecular geometry where atoms are arranged in a straight line, resulting in bond angles of 180 degrees. This arrangement often occurs in molecules with two central atoms or when hybridization leads to the formation of sp hybrid orbitals, which influences molecular shapes and bonding characteristics.
Nonpolar covalent bond: A nonpolar covalent bond is a type of chemical bond where two atoms share a pair of electrons equally, resulting in no significant charge difference across the bond. This occurs typically between atoms with similar electronegativities, meaning they have the same or very close attraction for electrons. Nonpolar covalent bonds are key for understanding molecular structures and behaviors, particularly in the context of molecular geometry and interactions.
Octet Rule: The octet rule is a chemical principle that states that atoms tend to bond in such a way that they each have eight electrons in their valence shell, achieving a stable electron configuration similar to that of noble gases. This rule is fundamental for understanding how and why atoms interact with each other through ionic and covalent bonding, as it drives the formation of stable compounds by promoting electron transfer or sharing among atoms.
Polar covalent bond: A polar covalent bond is a type of chemical bond that occurs when two atoms share electrons unequally, resulting in a molecule with a partial positive charge on one side and a partial negative charge on the other. This unequal sharing happens when the atoms involved have different electronegativities, meaning one atom has a stronger pull on the shared electrons than the other. The presence of a polar covalent bond leads to molecular polarity, which has significant effects on the physical and chemical properties of substances.
Resonance structures: Resonance structures are different Lewis structures for a molecule or ion that cannot be described by a single Lewis structure. These structures depict the same arrangement of atoms but involve different placements of electrons, particularly the locations of bonds and lone pairs. The actual molecule is a hybrid of all its resonance forms, illustrating the delocalization of electrons across bonds.
Single Bond: A single bond is a type of chemical bond that involves the sharing of one pair of electrons between two atoms. This bond is the simplest form of covalent bonding and is represented by a single line in Lewis structures, indicating that each atom contributes one electron to the shared pair. Single bonds are crucial in forming stable molecules and play a fundamental role in molecular geometry and reactivity.
Tetrahedral: Tetrahedral refers to a molecular geometry where a central atom is bonded to four other atoms, forming a shape resembling a tetrahedron. This configuration arises due to the spatial arrangement of the bonds, which minimizes repulsion between electron pairs around the central atom, leading to a three-dimensional shape that is important for understanding molecular interactions and properties.
Trigonal planar: Trigonal planar is a molecular geometry that describes a central atom bonded to three other atoms, with all bonds lying in a single plane and the bond angles measuring approximately 120 degrees. This arrangement results from the hybridization of orbitals, which leads to specific molecular shapes that minimize repulsion between electron pairs around the central atom.
Valence electrons: Valence electrons are the outermost electrons of an atom that are involved in forming chemical bonds. They play a crucial role in determining how atoms interact with each other, especially during the formation of covalent bonds and the creation of Lewis structures. The number of valence electrons varies among elements and helps explain trends in atomic properties across the periodic table.
VSEPR Theory: VSEPR Theory, or Valence Shell Electron Pair Repulsion Theory, is a model used to predict the three-dimensional shapes of molecules based on the idea that electron pairs around a central atom will arrange themselves to minimize repulsion between them. This theory connects the spatial arrangement of atoms within a molecule to the bonding and lone pairs of electrons present, which is essential for understanding molecular geometry, predicting angles, and visualizing hybridization.
Water (H2O): Water (H2O) is a simple molecule composed of two hydrogen atoms covalently bonded to one oxygen atom. This unique molecular structure gives water its distinctive properties, including its high specific heat, surface tension, and solvent capabilities, making it essential for life. The way hydrogen and oxygen share electrons leads to the formation of polar covalent bonds, which are critical in the context of biological systems and various chemical reactions.